The present invention relates to a method for recovering and/or removing metallic elements from waste water.
Environmental Hazard
Runoffs from a variety of industrial operations such as electrical power plants, steel plants, and mines are known to be contaminated with various metal compounds including iron, manganese, aluminum, zinc, copper, lead, arsenic, and chromium. These contaminants pose a serious environmental problem, as these runoffs cannot be safely discharged into the environment. Previously used methods to remove these contaminants involved adding lime, soda ash, or other neutralizing agents, and treating the runoff in a holding pond or clarifying tank. However, these methods have not been satisfactory because of the long periods of time required to effect treatment.
In particular, acid mine drainage from active and abandoned coal and metal mines is a serious environmental problem that affects thousands of miles of streams in the United States and elsewhere. Acid mine drainage, as well as mine tailings and refuse piles, also discolors streams. The metal sulfides that are part of many coal beds and ore deposits are oxidized when mining operations bring them within reach of oxidizing conditions created, directly or indirectly, by atmospheric free oxygen. This oxidation produces dissolved sulfate ion and metal ions. Because acidic wastewater, including acid mine drainage, discolors streams, damages ecological systems, and harms wildlife, federal and state limits for waste-water effluent require that the discharge pH be between 6 and 9, and that the concentration of common metallic elements such as iron and manganese be less than 2 milligrams per liter (mg/L) [cf. U.S. Environmental Protection Agency, 1994]. More stringent limits exist for toxic elements such as lead, arsenic, cadmium, and mercury.
Slow Oxidation by Oxygen and Extensive Ferric Iron Stains
Thermodynamically, aerated water should be capable of oxidizing ferrous iron (Fe2+) completely to ferric iron (Fe3+), with a equilibrium redox potential (also referred to as Eh when referenced to the standard hydrogen electrode, or ORP if not so specified) of:
O2+4H++4exe2x88x92=2H2O
Eh=1.23xe2x88x920.059 pH+0.015 log (pO2) voltsxe2x80x83xe2x80x83[Eq. 1]
This Eh value is shown as line [b] in FIG. 1, which is an Eh-pH diagram for the Fexe2x80x94H2O system.
At sea level, pO2=0.21 bars, and [Eq. 1 reduces to:
Eh=1.22xe2x88x920.059 pH volts [Eq. 1xe2x80x2].
In acid solutions, the ratio of ferrous ion and ferric ion concentrations is related to Eh as
Fe3++exe2x88x92=Fe2+
Eh=0.77+0.059 log(Fe3+/Fe2+) volts [Eq. 2]
By equating [Eq. 2] with [Eq. 1xe2x80x2], one can readily see that, at equilibrium, only a trace of Fe2+ should remain in solution, even when the pH is 3. When the pH increases, Fe3+ is precipitated as Fe(OH)3:
Fe3++3H2O=Fe(OH)3+3H+,
log(Fe3+)=4.84xe2x88x923 pH pH unitxe2x80x83xe2x80x83[Eq. 3]
and Fe3+ is precipitated as Fe(OH)3:
Fe(OH)3+3H++2exe2x88x92=Fe2++3H2O
Eh=1.06xe2x88x920.177 pHxe2x88x920.059 log(Fe2+) voltsxe2x80x83xe2x80x83[Eq. 4]
These equations show that, upon exposure to air or meteoric water saturated with air; iron should be precipitated as a solid phase, even in mildly acidic discharges, if true thermodynamic equilibrium were attained.
In reality, however, the Eh values of aerated waters are about 0.4 volts lower than the value calculated in [Eq.1xe2x80x2]. Sato [1960] measured the in situ Eh-pH values of mine waters at various depths, and found that, regardless of the type of deposits, the values in the oxidized zone were distributed within two parallel lines on an Eh-pH diagram, shown in FIG. 1. These lines are defined by the reaction:
O2+2H++2exe2x88x92=H2O2
Eh=0.68xe2x88x920.059 pH+0.0295 log[(pO2)/(H2O2)] voltsxe2x80x83xe2x80x83[Eq. 5]
The lower parallel line, marked [c] in FIG. 1, corresponds to the (pO2)/(H2O2) ratio of unity:
Eh=0.68xe2x88x920.059 pH voltsxe2x80x83xe2x80x83[Eq. 5xe2x80x2]
The upper parallel line, marked. [d] in FIG. 1, corresponds to the ratio of 106:
Eh=0.86xe2x88x920.059 pH voltsxe2x80x83xe2x80x83[Eq. 5xe2x80x3]
In laboratory oxidation of both Fe2+ and Mn2+ solutions, Sato [1960] showed clearly that the rate of Eh increase dropped drastically at the Eh value of [Eq. 5xe2x80x2] and never went over the value of [Eq. 5xe2x80x3] even after prolonged aeration. A plausible explanation is that O2 is somehow reluctant (i.e., a high activation energy barrier exists) to being split up in one step, and the faster reaction path is to form hydrogen peroxide as an intermediate. However, in the presence of iron, manganese, or a similar multivalent element, hydrogen peroxide is catalytically decomposed to oxygen and water, as discussed by Latimer [1952]. The result of this cyclic process is that oxygen becomes incapable of raising Eh beyond [Eq. 3xe2x80x3] in acidic solutions. Alternative paths may be provided by some aerobic microorganisms, but even with such help, the process of oxidation by oxygen is still relatively slow at surface temperatures.
The extremely slow rate of oxidation by oxygen beyond the Eh of [Eq. 3xe2x80x3] is the primary cause for the phenomenon of red iron stains formed for miles in the downstream direction from both treated and untreated acid mine discharge sites. When the pH is 2.5 to 4, ferric hydroxide can be precipitated by aeration alone, albeit slowly. The slow reaction ensures that a large fraction of the iron remains in solution as ferrous ion for a long time. Furthermore, the precipitation of iron as ferric hydroxide releases free sulfuric acid, which then becomes a secondary acidification step. Therefore, the pH often decreases downstream upon exposure to air even after neutralization treatment by anoxic limestone drains or by lime or caustic soda, which are bases used to bring the pH to more than 6 to meet discharge regulations.
The above pH range partially overlaps the range of active precipitation of gelatinous aluminum hydroxide (pH 4.5 to 6, Nordstrom and Ball [1986], and Hemingway [1982]; see FIG. 2) and aluminite, Al2(SO4) (OH)4.7H2O (pH 4.0, Robbins et al, [1996]). These poorly crystalline aluminum hydroxide and hydroxy-sulfate compounds and ferric hydroxide typically coat the limestone used in treatment, slowing down the neutralization process.
In FIG. 1, the range of pH was limited in this diagram to 1 to 11, because that is the pH range which is relevant to the present invention. The solid compound phases of iron are the most unstable hydroxide phases, Fe(OH)3 and Fe(OH )2, because these are the phases that actually precipitate upon addition of base, and also upon oxidation in the case of Fe(OH)2. The shaded area around these hydroxides is the region of active precipitation of iron as a solid from its ionic state in aqueous solution. The border on the ionic side is defined by the activity (concentration) of 10xe2x88x922 molar (560 mg/L Fe) and that on the solid side by the activity of 10xe2x88x925 molar (0.56 mg/L Fe) of Fe2+ or Fe3+ ion. The horizontal line that crosses the diagram at an Eh value of 770 mV indicates equal activities of ferrous and ferric ions. The four parallel lines with a slope of 59 mV per pH marked with a bracketed letter are as follows:
[a] the standard-hydrogen electrode i.e., H2xe2x80x94H2O redox couple;
[b] the standard potential of the H2Oxe2x80x94O2 couple;
[c] the standard potential of the O2xe2x80x94H2O2 couple; and
[d] the empirical limiting potential of oxidation by air.
Manganese and Sludge Problem
Manganese is more difficult to precipitate than iron when using air. The lowest pH that Mn2+ ion can be oxidized by air oxidation to a solid phase is about 6.5, if the Mn2+ concentration is as high as 550 mg/L, as shown in FIG. 3. To reduce the concentration to a more typical value of 0.55 mg/L without dilution, the minimum pH necessary for the oxidation is about 8.5. This is difficult to achieve with limestone alone, because the equilibrium of pure calcium carbonate at atmospheric partial pressure of carbon dioxide (3xc3x9710xe2x88x924 bars) is pH 8.3. Lime, slaked lime, caustic soda, or sodium carbonates are needed if the pH is less than 8.5. This alkalization not only consumes large quantities of basic reagents, it also produces equally massive waste material, i.e., sludge, which must eventually be disposed of in land fills.
In FIG. 3, the range of pH is given as 2 to 12, but this limit is merely for purposes of illustration. Manganese appears to precipitate as oxides, except for the purely Mn2+ state. The shaded area around these oxides is the region of active precipitation of manganese as a solid from its ionic state in aqueous solution. The border on the ionic side is defined by the activity of 10xe2x88x922 molar (550 mg/L Mn) and that on the solid side by the activity of 10xe2x88x925 molar (0.55 mg/L Mn) of Mn2+ or MnO4xe2x88x92 ion. The parallel lines marked with a bracket are as follows:
[a] the standard hydrogen electrode potential, i.e., H2xe2x80x94H2O redox couple;
[b] the standard potential of the H2Oxe2x80x94O2 couple;
[c] the standard potential of the O2xe2x80x94H2O2 couple; and
[d] the empirical limiting potential of oxidation by air;
[e] the estimated maximum redox potential for ozonated gas in [Eq4xe2x80x2].
Use of Ozone as and Oxidizer or Sanitizer
Ozone, O3, is a powerful and fast oxidizer, having a redox potential of:
O3+2H++2exe2x88x92=O2+H2O
Eh=2.08xe2x88x920.059 pH+0.0295 log(pO3/pO2) voltsxe2x80x83xe2x80x83[Eq. 6]
Ozone can be generated from pure oxygen or from air by passing the oxygen or air through a corona discharge field between two plates, or, more commonly, through concentric tubes of a dielectric material such as aluminosilicate glass, each backed by a thin metallic electrode. Alternatively, the oxygen or air can be passed through a silica glass tube irradiated by a ultra violet light source such as a mercury lamp. Ozone generators of various designs and capacities have been available commercially for at least 50 year.
The partial pressure of ozone in air that can be generated efficiently in modern devices is as much as 2 volume percent of oxygen, so that the practical upper limit of Eh of solution saturated with ozonated air is about:
xe2x80x83Eh=2.08xe2x88x920.059 pHxe2x88x920.0295 log(0.02)=2.03xe2x88x920.059 pHxe2x80x83xe2x80x83[Eq. 6xe2x80x2]
Ozone is used to disinfect and deodorize drinking water supplies in a large scale in many countries, including the United States, because it is converted to oxygen gas once oxidation is completed. Ozone is also used to degrade toxic organic and inorganic (e.g., cyanide) substances in ground water, sewage, and other waste water. More than 1600 papers were referenced by Rice [1984] regarding these and other applications of ozone. However, only a limited number of systematic studies of the use of ozone for removing metallic elements have been published. A few of these methods were directed to removing metallic elements. One example is the use of ozonated air to remove iron and manganese (0.1 to 0.5 mg/L level) and odor from pumped ground water that was contaminated by polluted, reducing water of the Rhine River [Weissenhorn, 1984].
A report specifically directed to acid mine drainage and not included in the bibliography of Rice [1984] was published by Rozelle and Swain [1974], who reported the results of static bench-top experiments that involved oxidation of near-neutral, pH-buffered Mn2+ solutions using ozone, hypochlorite ion, and chlorine gas. They found that ozone could reduce manganese solution concentrations by oxidizing Mn2+ from about 10 mg/L to 0.1 mg/L in relatively short times, about 1-5 minutes, at pH 7-8. These authors concluded that xe2x80x9cin order for ozonation to be useful for manganese removal from acid mine drainage, the acid mine drainage would have to be treated to remove iron because iron (II) is preferentially oxidized by ozone. This process would raise the pH. The ozonation would then be a secondary treatmentxe2x80x9d
Table 1 of this patent application show the redox potentials of elements that form solids from solution upon oxidation.
Use of Ozone for Recovery of Metallic Elements
To the best of the knowledge of the present inventors, there has been no literature published on this topic. A summary report made by Concurrent Technologies Corporation [1996] lists existing technologies to separate and/or purify recovered metal values. There is no disclosure of using ozone for this purpose.
Conventional treatments raise the pH of acidic waste water, such as the acid mine drainage, to 8.5 to 11 by adding large amounts of basic chemicals. After this, the contaminated water is impounded in large sedimentation ponds to allow atmospheric oxygen to oxidize and flocculate the metallic elements. This two-stage process produces large quantities of sludge, and is very slow and inefficient.
A number of other techniques have been disclosed to clean acidic waste water. Kinglsley et al., in U.S. Pat. No. 5,316,751, disclose using ozone to clean certain constituents of mine tail residues. In this process sands and slimes are separated from larger materials and the sands are separated from slimes. These slimes are deposited at a location at which the slimes are immersed in water, and the sands are deposited at a location where water can be drawn from the top surface of the slimes. The surface of the sands is sprayed with an aqueous solution of a leaching agent to leach metals and metal compounds from the sands. The slimes are injected with an aqueous solution of a leaching agent to leach metals and metal compounds from the slimes. Ozone can be used to break down resident cyanide before the residue is discharged. Metals and metal compounds are then recovered from the water drawn from the sands and slimes.
Yamasaki et al., in U.S. Pat. No. 5,580,458, disclose a method for treating waste water using a combination of calcium carbonate and microorganisms. Aluminum is added to remove fluorine from the wastewater.
Moniwa et al., in U.S. Pat. No. 5,492,633, disclose a process for treating water with ozone to oxidize and decompose trace amounts of organic substances contained in the water. A chelate compound is injected into the water to be treated after or just prior to the introduction of water into the reaction tank.
Murray et al, in U.S. Pat. No. 5,364,947, disclose a process for separating non-oxidizable compounds from a mixture containing at least one oxidizable compound by contacting the mixture with ozone to oxidize oxidizable compounds. These oxidized compounds are then converted to water-soluble hydrazones, which are then separated from the mixture using precipitation, liquid/liquid extraction, chromatography, etc.
Stevenson, in U.S. Pat. No. 5,370,800, discloses a method for removing metal compounds from waste water by adjusting the pH of the water to from 5 to 12, aerating the waste water, adding a flocculating agent to the water to flocculate metal compounds, and separating the flocculated metal compounds from the water.
Kazi et al., U.S. Pat. No. 4,752,412, disclose a method for recovering precious metals from ore by treating the ore in an acidic slurry with an activated oxygen mixture obtained from an ultraviolet light ozone generator. The activated oxygen frees chemically bonded precious metals, creating an expanded, hydrated ore so that the metals can be oxidized and leached out using standard leaching techniques.
Lindberg, U.S. Pat. No. 5,639,347, discloses a method for removing metals from acidic liquids containing dissolved metals by oxidizing the liquid with ozone, hydrogen peroxide, and/or air to increase the valence of the metals, making them easier to precipitate. After oxidation, the pH of the liquid is adjusted to over 6 to precipitate the metals.
Back et al., U.S. Pat. No. 5,607,653, disclose a process for treating and detoxifying salts-hydroxide scrubber wastes containing nitrite salts or sulfite salts. This process provides a continuous flow oxidation of the waste solution with ozone, followed by a neutralization step with monobasic potassium phosphate. Alternatively, oxidation and neutralization can be conducted in one step by adding hydrogen peroxide and monobasic potassium phosphate.
It is an object of the present invention to overcome the deficiencies in the prior art.
It is another object of the present invention to provide a process for recovering/removing metals from acidic waste water.
It is a further object of the present invention to provide a one-stage process for recovering/removing metals from acidic waste water.
It is another object of the present invention to provide a method for removing manganese stains from a substrate.
According to the present invention, ozone is used to rapidly oxidize specific metallic elements, including iron, manganese, lead, silver, nickel, cobalt, palladium, bismuth, thallium, and chromium which are present in acidic waste water either in a pond or in a flow reactor. The insoluble oxidized compounds of the metals formed by the ozonation are then recovered for industrial use in a conventional sedimentation/filtration tank or pool. There is no requirement for pre-treating or neutralizing the acidic waste water, even when iron is the dominant metal.
If the pH of the untreated acidic waste water is less than about 2.5, it is easy to separate iron from the other metals. After recovering the other metallic elements, iron is precipitated as ferric hydroxide using a conventional neutralizing agent such as finely powdered limestone.
Aluminum is removed as hydrated aluminum compounds by controlled neutralization after removal of the iron and prior to discharging the treated acidic waste water to streams.
Both the ozonation and neutralization processes are preferably monitored and controlled using electrochemical sensors and feedback controllers.
If the waste water has a near-neutral pH from pretreatment, but contains the above-mentioned metals in excess of allowed discharge limits, the process of the present invention can efficiently reduce the levels of these metals to allowable levels.